Chapter 8 Flashcards
Gases
- no fixed volume
- far more compressible than solids or liquids
- densities are very low (about 3-4 magnitudes less than solids and liquids)
- molecules are free to move over large distances
- pg160
The most important properties of a gas
Pressure, volume, and temperature
*pg160
Kinetic-molecular theory (4 points)
1) The molecules of gases are so small compared to the avg spacing bw them that the molecules themselves take up essentially no volume.
2) the molecules of gas are in constant motion, moving in straight line at constant speeds and in random direction bw collisions. The collisions of the molecules with the walls of the container define the pressure of the gas (the avg force exerted per unit area), and all collisions (molecs striking the walls and each other) are elastic (meaning the total KE is the same before and after the collision)
3) Since each molecule moves at a constant speed bw collisions and the collisions are elastic, the molecules of a gas experience no intermolecular forces.
4) The molecules of a gas span a distribution of speeds, and the avg KE of the molecules is directly proportional to the absolute temp (in K) of the sample: KE_avg ∝ T
* pg160
Ideal gas
A gas that satisfies all the requirements of the kinetic-molecular theory.
*pg160
Volume units and conversion
SI unit: cubic meter (m^3), but in chemistry cm^3 (or cc) and Liter (L) is more common.
1 cm^3 = 1cc = 1 mL and 1 m^3 = 1000L
*pg160
Temperature units
either in degrees Fahrenheit, degrees celsius, or in kelvins. The proper unit is K.
T (in K) = T(in ºC) + 273.15
*pg161
Pressure units
Defined as force per unit area. SI unit: pascal (Pa) where 1 Pa = 1 N/m^2.
At sea level, atmospheric pressure is about 101,300 Pa or 100 Kpa, this is 1 atmosphere (1 atm).
Torr is another common one where 1 atm = 760 torr. At 0ºC, 1 torr is equal to 1 mm Hg (millimeter of mercury) so generally 1 atm is equal to 760 mmHg.
1 atm = 760 torr = 760 mmHg = 101.3 kPa
*pg161
Standard temperature and pressure
Abbreviated STP. means a temp of 0ºC (273.15 K) and a pressure of 1 atm.
*pg161
Ideal Gas Law
The volume, pressure, and temp of an ideal gas are related an equation called the ideal gas law. This applies to most gas behaviour.
PV = nRT
P is the pressure; V is the volume; n is the #of moles of gas; R is the universal gas constant (0.0821 L-atm/K-mol) ; T is the absolute temp of the gas (in K)
equations can also be P1V1 = P2V2 (or rearranged not including the ones that are constant in both conditions)
*pg163
Boyle’s Law
If the volume decreases, the molecukes have less space to move around in. As a result they’ll collide with the walls of the container more often, and the pressure increases. If the volume of the container increases, the gas molecules have more available space and collide with the wall less often, resulting in a lower pressure.
P1V1 = P2V2
P ∝ 1/V
*pg163
Charles’s Law
If the pressure is to remain constant, then a gas will expand when heater and contract when cooled. if the temp of the gas is increased, the molecules will move faster, hitting the walls of the container with more force; in order to keep the pressure the same, the frequency of the collisions would need to be reduced, which is done by expanding the volume.
V1/T1 = V2/T2
V ∝ T
*pg163
Relationship between P and T from ideal gas law
If the temperature goes up, so does the pressure. As the temp increases the molecules move faster which means collide more often
P1/T1 = P2/T2
P ∝ T
*pg164
In a system with constant n:
summary of the relationships in ideal gas law
at constant P: V1/T1 = V2/T2 At constant T: P1V1 = P2V2 At constant V: P1/T1 = P2/T2 at only constant n: P1V1 / T1 = P2V2 / T2 *pg164
Avogadro’s Law
If two equal-volume containers hold gas at the same pressure and temperature, then they contain the same number of particles (regardless of the identity of the gas)
V1/n1 =V2/n2
V ∝ n
*pg165
Standard molar volume of an ideal gas at STP
the volume that one mole of a gas would occupy at 0ºC and 1 atm of pressure.
Equal to 22.4 L
*pg165
